Q13Engineering Chemistry

Electrochemical corrosion, universally known as 'wet corrosion', is the most common and destructive form of metal deterioration in everyday engineering. Unlike dry chemical corrosion which involves direct gas attack, wet corrosion absolutely requires the presence of a conducting liquid medium (an electrolyte, such as moisture, saltwater, or acidic soil). The fundamental mechanism relies entirely on the spontaneous formation of microscopic or macroscopic galvanic cells on the metal's surface.

The Galvanic Cell Mechanism

For wet corrosion to proceed, three essential components must be present: distinct Anodic areas on the metal, distinct Cathodic areas on the metal, and a continuous Electrolyte connecting them. These distinct areas can form due to impurities in the metal lattice, grain boundaries, mechanical stress, or variations in oxygen concentration (differential aeration).

• The Anodic Reaction (Destruction): The region that is relatively more electropositive or highly stressed acts as the Anode. It is exclusively at the anode where the actual physical destruction of the metal occurs. The solid metal atoms undergo vigorous oxidation, losing valence electrons and dissolving into the surrounding electrolyte as positively charged cations. General Equation: $M\to M^{n+}+n{e}^{-}$ Example for Iron: $Fe\to F{e}^{2+}+2e^{-}$
• Electron Flow: The electrons released by the dissolving metal cannot swim through the electrolyte. Instead, they instantly travel through the solid, conductive body of the metal itself toward the cathodic area.
• The Cathodic Reaction (Protection): The region that is relatively less active acts as the Cathode. The cathode itself does absolutely not dissolve or corrode. It merely acts as a conductive surface where the arriving electrons can be transferred to electron-accepting chemical species present in the surrounding electrolyte. The exact reduction reaction heavily depends on the pH of the environment: Case A - In Acidic Environments (like industrial atmospheres or acid rain): The dominant electron acceptors are Hydrogen ions ($H^{+}$). They gain electrons and evolve as hydrogen gas bubbles. $2H^{+}+2e^{-}\to H_2\uparrow$ Case B - In Neutral or Alkaline Environments (like normal seawater or soil): The dominant electron acceptor is dissolved Oxygen gas ($O_2$). Oxygen reacts with water and the incoming electrons to form Hydroxide ions ($O{H}^{-}$). $\frac{1}{2}{O}_2+H_{2}O+2e^{-}\to 2O{H}^{-}$

Ultimately, the positively charged metal ions migrating from the anode and the negatively charged hydroxide ions migrating from the cathode inevitably meet in the electrolyte and chemically combine to form the final visible corrosion product, such as Iron Hydroxide, which eventually dehydrates to form classic red rust ($F{e}_2{O}_3\cdot x{H}_{2}O$).